Chemistry

Key unit competency: Compare and contrast the properties of the Period 3 elements and their compounds in relation to their positions in the Periodic Table.

Learning objectives:

By the end of this unit I will be able to:

•Compare the physical properties of the Period 3 elements.

•Describe the nature of the oxides of the Period 3 elements and the type of bonding in their chlorides, oxides and hydrides.

•Relate the physical properties of the Period 3 elements to their position in Periodic Table.

•Relate the physical properties of compounds of the Period 3 elements to their nature of bonds across the period.

14.1. Physical Properties of the Period 3 elements

Activity 14.1

1. Write the electronic configuration of the following elements in terms of s, p, d and f...(i) Sodium (ii) Magnesium (iii) Aluminium (iv) phosphorous (v) sulphur

2.Considering the electronic configuration of magnesium and Aluminium, phos-phorus and sulphur. How do you expect their ionization energies to vary?

3. How do you expect the general trend in ionization energy, electron affinity, melting and boiling point, electronegativity to vary for the elements in the pe-riod 3?

4. Considering the electronic configuration of magnesium and Aluminium, phos-phorus and sulphur. What can you say about them, how do you expect their ionization energies to vary?

(a) Variation of First ionization energies (IE) of Period 3 elements

First ionization energy generally increases across Period 3 from left to right. However, it drops at aluminium and Sulphur (table 14.1 and Fig.14.1). This can be exaplained in term of more stable electronic structures of the two elements after loosing 1 electron:

Table 14.1: Variation of first ionization energies of period 3

Going across Period 3, there are more protons in each nucleus so the nuclear charge in each element increases. Therefore the force of attraction between the nucleus and outer electron is increased, and there is a negligible increase in shielding because each successive electron enters the same energy level. So apart from the two exceptions mentioned above, the first ionization energy increases from left to right in the period (figure 14.1).

b. Variation of atomic radius of Period 3 elements

Table 14.2: Variation of atomic radius of period 3 elements

Going across Period 3 from left to right, the number of protons in the nucleus increases so, the nuclear charge increases. There are more electrons, but the increase in shielding is negligible because each extra electron enters the same principal energy level. Therefore, the force of attraction between the nucleus and the electrons increases. So the atomic radius decreases as indicated in the Figure 14.2 and table 14.2.

(c)Variation of electronegativity of Period 3 elements

Table 14.3: Variation of electronegativity of period 3 elements

Going across Period 3 from left to right, electronegativity increases almost linearly due to the nuclear charge increase as atomic radius decreases. There are more electrons, but the increase in shielding is negligible because each extra electron enters the same principal energy level so electrons will be more strongly attracted to the nucleus.

You might expect argon (with 18 electrons) to be the most electronegative element in Period 3, but its outer energy levels are full. Therefore, it does not form covalent bonds with other atoms, so it is given an electronegativity value of zero.

d. Variation of melting and boiling points in Period 3

Melting and boiling points generally increase going from sodium to silicon, then decrease going to argon with a “jump” at Sulphur (Fig 14.4 and Table 14.4).

Table 14.4: Variation of melting and boiling points of period 3 elements

Sodium, magnesium and aluminium are all metals. They have metallic bonding, in which positive metal ions are attracted to delocalized electrons. Going from sodium to aluminium, the charge on the metal ions increases from +1 to +3 through magnesium at +2, the number of delocalized electrons increases, so the strength of the metallic bonding increases and the melting points and boiling points increase.

Silicon has giant covalent bonding structure. It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement. This extends in three dimensions to form a giant macromolecule and this explains the very high melting point and boiling point.

Phosphorus, sulphur, chlorine and argon, are all non-metals, and they exist as small, separate molecules. Phosphorus, sulphur and chlorine exist as simple molecules. Argon exists as separate atoms (monatomic). When these these elements melt or boil, it is the van der Waals’ forces between the molecules which are broken; being weak they need little energy to overcome them.This explains why their melting and boiling points are low. However, Sulphur has a higher melting point and boiling point than the other three because

Phosphorus exists as P4 molecules, Sulphur exists as S8 molecules, chlorine exists as Cl2 molecules, and argon exists as individual Ar atoms.

The strength of the van der Waals’ forces decreases as the size of the molecule decreases, so the melting points and boiling points decrease in the order S8> P4> Cl2>Ar

e. Variation of electrical conductivity of Period 3 elements

Electrical conductivity increases going across Period 3, left to right, from sodium to aluminium, then decreases sharply to silicon as indicated by the graph 14.5 below.

Table 14.5: Variation of relative electrical conductivity of period 3 elements

The delocalized electrons are free to move and carry charge. Going from sodium to aluminium, the number of delocalized electrons increases, there are more electrons which can move and carry charge so the electrical conductivity increases.

Silicon is called a semi-conductor because at higher temperatures more electrons are promoted to the higher energy levels so there are more delocalized electrons to move and carry charge.

Phosphorus, sulphur and chlorine, the outer electrons are not free to move and carry charge because they are held strongly in covalent bonds. In argon (mono atomic) the outer electrons are not free to move and carry charge because they are held strongly in a stable third energy level and this explains their zero electrical conductivity.

f. Variation of metallic character of period 3 elements

Metallic character decreases as you move across a period 3 in the periodic table from left to right. This occurs as atoms more readily accept electrons to fill the valence shell than lose them. Note that as the metallic character decreases across the period, the reducing power decreases whereas oxidizing power increases.

g. Variation of electron affinity across period 3 elements

The electron affinity [EA] is the energy change for the process of adding an electron to a gaseous atom to form an anion (negative ion).

As we might predict, it becomes easier to add an electron to an atom across the period from left to right as the effective nuclear charge of the atoms increases. As we go from left to right across a period 3, EAs tend to become more negative, i.e. ability to acquire electrons increases.

The exceptions found among the elements of group 2 (2A), group 15 (5A), and group 18 (8A) can be understood based on the electronic structure of these groups.

Magnesium and phosphorous have anomalous electron affinity, Magnesium has a positive EA while phosphorous is less negative. This is due their electron arrangement, where in magnesium the electron should be added to 2p orbital which is a less stable structure compared to the electronic structure of the atom [Ne]3s2. Similarly in phosphorous the electron should be added to 3p orbital which is half filled and thermodynamically stable.

Checking up 14.1

1.Explain the variation of the following terms as applied in period 3 of the periodic table:

(i) Ionization energy, Electronegativity, )

(ii)Explain the anomalous behavior indicated by magnesium and phospho-rous in graph 14.1 above

2.The table below shows the melting points of the period 3 elements except for silicon:

(a)Explain in terms of bonding why the melting point of magnesium is higher than that of sodium.

(b) Predict the approximate melting point of silicon.

(c) Explain why chlorine has a lower melting point than sulphur.

(d) Explain the variation of metallic character, electronegativity, atomic radii ,first ionization energy, melting and boiling points, electron affinity and electrical conductivity across the period

14.2. Chemical properties of period 3 elements

Activity 14.2

a.Experiment to investigate the action of water on period 3 elements

Materials /apparatus

Water , test tubes, a piece of sodium metal, aluminium powder, magnesium ribbon/powder, red litmus paper/ universal indicator

Procedure:

1. Put about 5cm3 of water in each of the three test tubes arranged in a test tube rack

2. Add a small piece of Mg, Na and Al in each of the three test tubes in 1.

3. Record your observations in a suitable table.

Study questions:

1. What do you say about your observations made in experiment above.

2. Write equation for the reaction that occurs in each test tube in procedure 2.

(b) Experiment to investigate the action of heat on period 3 elements

Materials /apparatus:

Water , test tubes, a piece of sodium metal, aluminium power/sheet, magnesium ribbon/powder, phosphorous and sulphur powder, universal indicator , pair of tongs, source of heat

Procedure:

1. Hold a piece of magnesium ribbon on a Bunsen flame and record you observa-tion.

2. Repeat experiment 1 for sodium, aluminium, phosphorous and sulphur and record your observation in each case.

3. For each of the products formed i.e. for metal oxides formed, add water and dip a litmus paper to test their nature.

Note: if the oxide is gaseous hold a piece of litmus paper on the mouth of the test tube.

Study questions:

1.Write equations to show how the metals react with oxygen.

2. What would you expect to observe when the metal is burned in oxygen.

a) Reaction with water

Reactivity with water generally decreases across the period from left to right because there is a decrease in metallic properties.

i) Sodium reacts vigorously with cold water to form sodium hydroxide and hydrogen gas.

ii) Magnesium does react slowly with cold water to form magnesium hydroxide solu-tion, but with steam the reaction is faster to form magnesium oxide and hydrogen gas.

iii) Aluminium reacts with steam very slowly to form aluminium hydroxide and hy-drogen gas.

Aluminum powder heated in water vapor produces hydrogen and aluminum oxide. The reaction is relatively slow because of the existing strong aluminum oxide layer on the metal, and the build-up of more oxide during the reaction.

iv)Silicon reacts a little bit faster with steam at red hot to produce silicon dioxide and hydrogen.

v. Phosphorus and sulfurhave no reaction with water

.vi)Chlorine dissolves in water to give a green solution. A reversible reaction produces a mixture of hydrochloric acid and chloric(I) acid (hypochlorous acid).

In the presence of sunlight, the hypochlorous acid slowly decomposes to produce more hydrochloric acid, releasing oxygen gas:

There is no reaction between of argon and water that is known.

b. Reaction with oxygen

Elements must be heated to react with oxygen; however, dry white phosphorus can ignite spontaneously and that is why it is stored under water. The reactivity depends much on the state of subdivision

Sodium and magnesium: Sodium and magnesium burnvigorously in oxygen to form ionic white oxides

Aluminium: sheets of aluminium get slowly coated with thin oxide layer of aluminium oxide.

c. Reaction with NaOH

d. Reaction with HCl.

e. Reaction with Hydrogen

f. Reducing and oxidizing power

Elements on the left of the period three are metal.They react by losing their valance electrons; hence they are good reducing agents.Their reducing power decreases from Na to Al.

The elements on the right of the period three are nometals. They react by gaining or sharing their valance electrons; they are good oxidising agents.Their oxidising power increases from Si to Cl.

In general the oxidizing power increases from left to right and the reducing power decreases from left to right across period 3 due to decrease of atomic size that affects the ionization energy and electronegativity.

Checking up 14.2

1. Describe the nature of hydrogen compounds of period 3.

2. Explain why the reducing power of period 3 elements decreases across the period.

14.3. Compound of period 3 elements

The oxides of period 3 elements:

Activity 14.3(a)

1. Write the formulae of the oxides of period 3 elements

2. What did you consider when writing the formulae of the oxide in 1 above?

3. How do you expect the oxides to behave in water? Explain your answer.4. Suggest the trend of acid- base character of the oxides of period 3

The metallic oxides on the left of the period adopt giant structures of ions. In the middle, silicon forms a giant covalent oxide (silicon dioxide); the elements on the right form simple molecular oxides with simple structures. The intermolecular forces binding one molecule to its neighbors are van der Waals dispersion forces or dipole-dipole interactions.

Physical properties of the oxides of period 3 elements

Melting and boiling points: the metal oxides and silicon dioxide have high melting and boiling points because a large amount of energy is needed to break the strong bonds (ionic or covalent) operating in three dimensions. The oxides of phosphorus, sulfur and chlorine consist of individual molecules.

Electrical conductivity: None of the oxides above have any free or mobile electrons, indicating that none of them will conduct electricity when solid.

Phosphorus has two common oxides, phosphorus (III) oxide, P4O6, and phosphorus(V) oxide, P4O10.

Phosphorus (III) oxide is a white solid, melting at 24°C and boiling at 173°C.

Phosphorus(V) oxide is also a white solid, which sublimes at 300°C. In this case, the phosphorus uses all five of its outer electrons in the bonding.

Sulphur has two common oxides, sulphur dioxide (sulphur(IV) oxide), SO2, and

sulphur trioxide (sulphur(VI) oxide), SO3.

Sulphur dioxide is a colorless gas at room temperature with an easily recognized pungent smell. It consists of simple SO2 molecules.

Pure sulphur trioxide is a white solid with a low melting and boiling point. It reacts very rapidly with water vapor in the air to form sulfuric acid. Under laboratory

conditions, it forms a white sludge which fumes dramatically in moist air, forming a fog of sulfuric acid droplets.

Chlorine forms several oxides. Two are considered here: chlorine(I) oxide, Cl2O, and

chlorine(VII) oxide, Cl2O7.

Chlorine (I) oxide, Cl2O, is a yellowish-red gas at room temperature. It consists of simple, small molecules.In chlorine (VII) oxide, Cl2O7, the chlorine losesall of its seven valence electrons in bonds with oxygen. This produces a molecule much larger than chlorine (I) oxide, suggesting higher melting and boiling points. Chlorine (VII) oxide is a colorless oily liquid at room temperature.

Checking up 14.3(a)

1. The table below shows oxides of period 3 in the periodic table

(a)Copy and complete the table using the following guidelines

(i)Complete the ‘bonding’ row using only the words ionic or covalent

(ii) Complete the ‘structure’ row using simple molecular,giant ot lattice

(iii) Explain in terms of forces, the difference between the melting points of MgO and SO3.

b. The oxides Na2O, Al2O3, SO3, were each separately added to water. For each oxide, construct a balanced equation for its reaction with water

Acid-base Behavior of the Oxides

Activity 14.3 (b)

1. Classify the oxides in terms physical states of the oxides of period 3.

2. How do you expect the oxides react with water, acids, and sodium hydroxide.(use equations to justify your answer)

3.(a) Predict the nature of oxides of period 3 elements when dissolved in water.

(b)What would you expect to observe when both blue and red litmus papers are dropped into each of the solutions formed in question (2) above in water.

Acidity increases from left to right, ranging from strongly basic oxides on the left to acidic ones on the right, with an amphoteric oxide (aluminum oxide) in the middle.

Checking up 14.3(b)

1.Consider the following oxides: CaO, Al2O3 , Na2O , MgO , P2O5 , SO2 SiO2.

(a)State which of the oxides arei)Basic (ii) acidic (iii) neutral (iv)Amphoteric

(b)Describe briefly how you could test to see whether a solid oxide is basic , acidic , neutral or Amphoteric 2. (a)In each of the following species state whether the metal to non metal bond is ionic or covalent and explain why : Na2O , Al2O3

(b) An uknown inorganic oxide is a white solid of melting temperature 17100C and is uncreative towards water.(i)State and explain whether or not this is sufficient information to deduce whether the bonding is ionic or covalent.(ii) Outline one additional test which could be applied to support your conclusion.

14.3.2. Chlorides of Period 3 Elements

Activity 14.3(c)

The table below gives some data about the chlorides of elements of period 3

(a) Explain why the boiling point of NaCl is higher than that of MgCl2.

(b) Make a research and explain why the pH of NaCl is 7 and that of AlCl3 is 2.

The chlorides of interest are given in the table below:

Sulphur forms three chlorides, but S2Cl2 is the most common. Aluminum chloride may be found, in certain conditions, as a dimmer, Al2Cl6. Phosphorous forms two chlorides, PCl3 and PCl5.

The nature and structure: Sodium chloride and magnesium chloride are ionic and consist of giant ionic lattices at room temperature. Aluminum chloride exhibits covalency characters. Aluminium ion has high charge density; due to this, the electron cloud of the chloride is distorted toward the aliminium ion, Al3+, impacting an appreciable covalent property to the bond.

The other chlorides are simple covalent molecules.

Melting and boiling points of Period 3 chloride

Sodium and magnesium chlorides are solids with high melting and boiling points because of the large amount of heat which is needed to break the strong ionic attractions.

Aluminum chloride and phosphorus(V) chloride are solids with relatively low melting and boiling points. The remaining chlorides are volatile liquids due to the weak van der Waals that hold their molecules together.

Electrical conductivity: solid chlorides do not conduct electricity because the ions are not free to move

Sodium, magnesium and aluminium chlorides are ionic and so will conduct electricity when they are molten or in aqueous solution. The rest of the chlorides do not conduct either in solution or molten state due to absence of ions.

Reactions with water

The chlorides of more electropositive metals (NaCl and MgCl2) dissolve in water to form neutral solutions. The other chlorides (covalent) hydrolyze or react with water to form acidic solutions. As shown below:

The greater positive charge on aluminium attracts electrons in the water molecules quite strongly, making O-H bond of water molecule in the complex ion so weak to dissociate and liberate H+ ion.

Checking up 14.3(c)

1. a. Distinguish between dissolving and hydrolysis.

b. Name one chloride that dissolves in water, and one chloride that undergo hydrolysis.

c. State how the bonding in the chlorides changes on crossing the second and third periods from left to right

d. Suggest two ways how you would know that a reaction has taken place when a few drops of water is added to silicon (iv) chloride.

2. a. Write an equation that shows thata solution of beryllium chloride in water is acidic.

b. Predict the shape of PCl3 molecule and suggest a likely bond angle.

c. Iron (iii) chloride, FeCl3, forms dimmers in the gas phase similar to those of aluminium chloride. Draw the likely structure of these two dimers.

4. Carbon and silicon are in the same group of the periodic table and forms chloride of CCl4 and SiCl4 respectively. However, CCl4 does not undergo hydrolysis whereas SiCl4 hydrolyses in water to form white fumes of the gas.

14.3.3. The hydrides of period 3 elements

Activity 14.4(d)

1. Period 3 elements from sodium to chlorine form different hydrides of different bond nature, physical properties and structure.

(a) Write the formula of the hydrides formed by period 3 elements.

(b) Predict the nature of bonding based on your knowledge of periodicity of ele-ments in the periodic table.

(c ) Basing on the nature of bonding predicted in (b) above. How would you ex-pect their boiling and melting point vary across the period?

(d) Predict the nature of solutions formed by hydrides when dissolved in water. What would you expect to observe if red and blue litmus papers were separate-ly dropped into each solution?

Hydrides are commonly named after binary compounds that hydrogen forms with other elements of the periodic table. Hydride compounds in general form with almost any element, except a few noble gases. The common hydrides of period 3 elements are as shown in the table 14.5 below

The hydrides above are examples of period 3 elements with some of their properties summarized in the tabl. As we can see the hydrides of period 3 vary from ionic hydride such as NaH at the left side to polar covalent hydride such as HCl at the right side of the period.

Checking up 14.3(d)

1. The table below gives some properties of oxides of period 3 in the periodic table.

a. Explain the trend in the melting points of the above oxides and their origin.

b. Write the equations to show the reaction of

i. SO3 with water

ii. Na2O

c. One of the above oxides has acidic and basic character.i. Identify that oxide from the above table ii. Using equations explain the acidic and basic character of the oxide identified in (c) (i).

2. This question concerns the following oxides: Na2O, MgO, SiO2, and SO3.From the list above identify the oxide that best fits the description given:

i. An oxide that is insoluble in water.

ii. An oxide that has simple molecular structure at room temperature and pressure.

iii. An oxide that reacts with water forming a strongly alkaline solution.

iv. An oxide that is slightly soluble in water forming weak alkaline solution.

3. There is a link between the properties of the oxides of the Period 3 elements and their structure and bonding. The table below shows the melting points of the oxides of some Period 3 elements.

a. In terms of crystal structure and bonding, explain in each case why the melting points of sodium oxide and silicon dioxide are high.

b. Predict whether the melting point of lithium oxide is higher than, the same as, or lower than the melting point of sodium oxide and explain your prediction.

c. Phosphorus (V) oxide has a lower melting point than sodium oxide.

i. State the structure of and bonding in phosphorus (V) oxide.

ii. Explain why the melting point of phosphorus(V) oxide is low.

d. Samples of phosphorus(V) oxide and sodium oxide were reacted with water. In each case, predict the pH of the solution formed and write an equation for the reaction.

4. Sodium chloride is a high melting point solid which dissolves in water to make a colorless solution. Silicon (IV) chloride is a liquid at room temperature which fumes in moist air, and reacts violently with water.

a. Draw a diagram to show the arrangement of the particles in solid sodium chloride, making clear exactly what particles you are talking about.

b. Explain why this arrangement leads to a high melting point.

c. Draw a simple diagram to show the structure of silicon (IV) chloride, and explain why silicon (IV) chloride is a liquid at room temperature.

d. Why is there such a big difference between the chlorides of sodium and silicon?

e. Briefly describe and explain the difference in electrical conductivity between sodium chloride and silicon (IV) chloride in both solid and aqueous molten state.

f. Write an equation to show what happens when silicon (IV) chloride reacts with water.g. Name another Period 3 chloride which behaves similarly to sodium chloride, and one which behaves similarly to silicon (IV) chloride.

5. With the help of equation describe how the hydrides of period 3 react with water.

14.5. End unit assessment

1. Use the information in the following table to explain the statements below

(a)What is the trend in atomic radius? Explain the origin of that trend.

(b) The ionic radii of Na+, Mg2+ and Al3+ are less than their respective atomic radii , whereas the ionic radii of Cl- and S2- are greater than their respective atomic radii.Compare the atomic radii with the ionic radii and explain what your observation

(c) What trend do you observe in the 1st ionization energy? Explain the origin of that trend?

d. The first ionization energy of Al is less than that for Mg; why?

2. The table below shows the melting points of the period 3 elements except for silicon.

a. Explain in terms of bonding why the melting point of magnesium is higher than that of sodium

b. State the type of bonding between atoms in the element silicon and name the type of structure which silicon forms.

c. Predict the approximate melting point of silicon.

d. Explain why chlorine has a lower melting point than sulphur.

e. Predict the approximate melting point of potassium and give one reason why it is different from that of sodium.

3. The elements Sodium, Magnesium, silicon, phosphorous and chlorine are members of the third period of the periodic tablea.

a.i. Write down the formula of the principal oxides and chlorides of the elements listed above and in each case indicate the type of bonding.

ii. Explain what happens when each of the above oxides and chloride is added to water and indicate whether the resultant solution will be acidic, basic or neutral.

c. The melting points of Mg , Si and S are 6500C, 14230C respectively. Explain the differences in the melting points of the elements.

d. Name the type bonding that exists in the hydrides of the elements Sodium, Phosphorous and sulphur and write the equations to show the reactions if any of the hydrides with water.

4. Choose from the elements: Sodium, magnesium aluminium, silicon, phospho-rous, chorine and argon a. List the elements that react readily with cold water to form alkaline solutions. And write the equations for the reactions.

b. List the hydrides that have hydrides with low boiling points/temperatures and explain why.

c. List the elements that form nitrates and write the formulae of nitrates.

d. What is the most ionic compound that can be formed by the combination of two of these elements.

e. Which element has both metallic and non metallic properties?

f. Name the elements that normally exist as molecules.

5.a. Describe the formation of Al2Cl6 dimers from AlCl3 monomers in terms of or-bital hybridization

b. i. Which of the oxides listed above form oxide of the type X2O3

ii.Describe briefly how you would prepare each of the oxide and give the equations for the reactions involved.iii. What would be observed if each of the above oxide in (b)

(ii) was reacted with sodium hydroxide? Write equation for the reactions involved in each case